SCIENCE 1206 - Weebly
SCIENCE 1206
UNIT II: CHEMICAL REACTIONS Text: Chapters 5-8
Chemistry: the study of matter, its properties and its changes
Matter: anything that has mass and takes up space (energy is not matter).
The three states of matter: solid, liquid and gas
Classification of Matter as Pure Substances or Mixtures:
Pure Substances: have constant composition; all the particles that make up the substance are the same
1. Elements: - the simplest form of matter that can exist under normal conditions
- composed of only one kind of atom eg. C, Mg
- cannot be broken into simpler substances by chemical means (heat/electricity)
- combine to form other substances
2. Compounds: - substances composed of two or more different kinds of atoms eg. H2O, NaCl
- can be broken down into simpler substances by chemical means
Mixtures: have variable composition: composed of 2 or more pure substances
1. Homogeneous Mixtures: solutions – have only one visible component
eg. tap water, air, sugar solution (sugar + water)
2. Heterogeneous Mixtures: mechanical mixtures – have 2 or more visible components
eg. sand in water, vegetable soup
Properties and Changes of Matter:
1. Physical Property: characteristics of matter, used to identify substances
eg. state at room temperature, boiling and melting points, color, solubility,
mass, electrical conductivity
2. Physical Change: a change in the size or form of a substance that does not change its composition
eg. cutting, bending, changes in state: boiling, melting, condensing, and solidifying
3. Chemical Property: characteristic of matter that can be observed when matter undergoes a
change in composition (chemical reaction): describes "how it reacts"
eg. butane reacts with oxygen to produce carbon dioxide and water
4. Chemical Change: a chemical reaction; a change in which at least one or more new substances
(products) are formed. The products have different properties from the starting substances (reactants).
eg. Fe(s) + O2(g) ( Fe2O3(s) The rust produced has completely different
properties from iron and oxygen.
Evidence of Chemical Change: change in color, odor, energy (temperature change, light)
change in state: bubbles = new gas produced
precipitate = new solid produced
ELEMENTS & THE PERIODIC TABLE
All elements are classified as metals or nonmetals, depending on their properties.
| | | |
|PROPERTY |METALS |NONMETALS |
| | | |
|LUSTRE |shiny |dull |
| | | |
|MALLEABILITY |malleable (bendable) |brittle |
| | | |
|CONDUCTIVITY OF HEAT & ELECTRICITY |good conductors |poor or nonconductors |
| | | |
|STATE AT ROOM TEMPERATURE |all solids except mercury, |most are gases, some are solids |
| |Hg = liquid |and bromine, Br = liquid |
| | | |
|REACTIVITY WITH ACID |mostly yes |no |
| | | |
|LOCATION (PERIODIC TABLE) |left of staircase line |right of staircase line |
METALLOIDS (Semimetals)
- elements that have some properties of metals and some properties of nonmetals
- include all elements on either side of the staircase line except Al and At
- also includes one form of Carbon, graphite, which is dull and brittle (nonmetal),
but is a good conductor of electricity (metal)
CHEMICAL FAMILIES (GROUPS):
Groups of elements in the same vertical column that have similar physical and chemical properties.
1. Alkali Metals: - Group 1, IA
- show metallic properties (see table above)
- highly reactive, especially with water; reactivity increases going down the group
- Cs & Fr are the most reactive metals
2. Alkaline Earth Metals: - Group 2, IIA
- show metallic properties (see table above)
- less reactive than alkali metals; reactivity increases going down the
group.
Note: Metals from both group 1 and group 2 react with water to form alkaline (basic) solutions.
3. Halogens: - Group 17, VIIA
- show nonmetallic properties (see table above)
- reactivity decreases going down the group: F is the most reactive nonmetal
- react with most metals to produce salts (ionic compounds)
- react with hydrogen to form compounds that dissolve in water to form acids
4. Noble Gases: - Group 18, VIIIA
- show nonmetallic properties
- extremely low chemical reactivity
SERIES OF ELEMENTS:
1. Representative Elements: A groups or groups 1, 2, 13 –18
2. Transition Elements: B groups or groups 3 – 12
HYDROGEN: - the lightest element and most abundant element in the universe
- doesn't really belong to any group
- it sometimes behaves like an alkali metal, sometimes like a halogen and at other times in
its own unique way ie. as an acid
PERIODS: horizontal rows of the periodic table
THE ATOM:
• The basic building block of all matter
• The smallest particle of an element that retains the properties of that element
• Electrically neutral: # of positive charges = # of negative charges
• composed of 3 types of subatomic particles:
| | | | | |
|PARTICLE |SYMBOL |RELATIVE CHARGE |ACTUAL |LOCATION |
| | | |MASS (g) | |
| | | | | |
|Proton |p+ |1+ |1.67 x 10-24 |nucleus |
| | | | | |
|Neutron |no |0 |1.67 x 10-24 |nucleus |
| | | | | |
|Electron |e- |1- |9.11 x 10-28 |orbital |
Atomic Number: - identifies the element
- equal to the number of protons in the nucleus
- since atoms are electrically neutral, # of protons = # of electrons
Mass Number: - # of protons + # of neutrons
- protons and neutrons account for most of the mass of the atom
Quantum Mechanics Theory of the Atom:
According to this theory, an electron with a specific energy occupies a region in space (orbital) or
electron energy level.
Electron Energy Diagrams of Atoms:
• An energy level represents a specific value of energy of an electron and corresponds to a general location
• The number of occupied energy levels in any atom is normally the same as the period number in which the atom appears
• for the first 3 energy levels, the maximum number of electrons that can be present are 2, 8 and 8 in order of increasing energy (increasing distance from nucleus)
• a lower energy level is filled with electrons to its maximum before the next level is started.
• the electrons in the highest (outermost) occupied energy level = valence electrons, which is the same as the group number (for group A elements)
Assign: Worksheet #1, p. 12
STABLE ATOMS
• have low chemical reactivity
• include noble gases, all of which have 8 valence electrons (except He, which has 2)
• other atoms can become more stable by reacting and changing the number of their electrons, thereby attaining the same stable electron configuration (structure) of the nearest noble gas:
• atoms can follow one of two rules:
a) Octet Rule: - atoms attempt to obtain 8 valence electrons
- includes most atoms
b) Duet Rule: - atoms attempt to obtain 2 valence electrons
- includes H, Li and Be
• one way atoms can achieve a stable octet or duet is by forming ions
ION
• an atom or group of atoms that has a positive or negative charge, due to the loss or gain of one or more electrons
• single atoms: form simple ions (monatomic ions)
• group of atoms: form complex ions (polyatomic ions)
Example: Sodium metal and chlorine gas react to produce NaCl, a very stable and unreactive substance, compared to Na (alkali metal) or Cl (halogen).
They do so by first forming ions.
_______
_______ _______ _______
_______ _______ _______
Compare to nearest
Noble gas:
Na atom Na+ ion
_______ _______ _______
_______ _______ _______
_______ _______ _______
Compare to nearest
Noble gas:
Cl atom Cl- ion
Summary: When sodium metal and chlorine gas react, the sodium atoms each lose one electron to a chlorine atom. In so doing the atoms form ions of opposite charge:
Cations: - Ions which form when atoms lose one or more valence electrons
- metal atoms form cations
- for group A atoms, # of valence electrons lost = group # = charge on ion
- have a positive charge because they have more protons than electrons
Anions: - Ions which form when atoms gain one or more valence electrons
- nonmetal atoms form anions
- for group A atoms, # of valence electrons gained = 8 – group # = charge on ion
- have a negative charge because they have more electrons than protons
Naming Ions:
Cations: element name + the word “ion” eg: Na+ = sodium ion
Anions: stem of element name + ide + the word “ion”
eg: Cl, chlorine becomes Cl- = chloride ion
P, phosphorus becomes P3- = phosphide ion
O, oxygen becomes O2- = oxide ion
Note:
1. Both cations and anions are more stable than the atoms from which they form since these ions
attain the same stable electron configuration as the nearest noble gas.
2. Boron, carbon and silicon do not tend to form ions (they instead share electrons with other atoms)
3. The noble gases do not form ions since they are already stable (have filled orbitals)
4. Hydrogen can form a cation or an anion:
• Cation: H+ , hydrogen ion has 1 proton but no electrons
• Anion: H- , hydride ion has 1 proton and 2 electrons
Assign: Worksheet #2, p.13 and Worksheet #3, p.14
IONIC COMPOUNDS
• All are solids at SATP (Standard Ambient Temperature and Pressure) of 25oC and 100 kPa.
• When they dissolve in water, they form aqueous solutions that:
• are colored or colorless
• conduct electricity ie. they are electrolytes
• These compounds form after an electron transfer:
• usually from a metal to a nonmetal
• the resulting ions (cations and anions) are attracted to each other (since they are oppositely charged) and they form ionic bonds
• Together all of the ions present form an ionic crystal lattice in which the net charge is zero
• Eg: (1) in a sample of sodium chloride, NaCl, for every Na+ ion there is one Cl- ion
• Eg: (2) in a sample of calclium chloride, CaCl2, for every Ca2+ ion there are 2 Cl- ions
• Formula Unit: an expression of the simplest whole number ratio of cations to anions
Eg: NaCl 1:1 Ratio of Na+ : Cl-
CaCl2 1:2 Ratio of Ca2+ : Cl-
Types of Ions and Ionic Compounds
A. Monatomic Ions (Simple Ions)
• Single atoms that have lost or gained one or more electrons
• Form Binary Ionic Compounds (2 simple ions)
• Eg. Na+ Cl-
B. Polyatomic Ions (Complex Ions)
• Cations or anions composed of a group of atoms with a net positive or negative charge
• Eg. NH4+ NO2- NO3- CO32-
Ammonium ion Nitrite ion Nitrate ion Carbonate ion
C. Multivalent Ions
• certain transition metals can form more than one type of ion, each with a different charge
• Eg. Fe3+ Fe2+
• The more commonly occurring is listed on top, thus Fe3+ is more common than Fe2+
D. Hydrated Ionic Compounds
• Water molecules are loosely held within the ionic compound
• Eg. ZnCl2 • 6H2O CuSO4 • 5H2O
Writing Empirical Formulas
• Chemical formulas for ionic compounds
General Rules
1. Write each ion symbol with charge: simple ions from front of periodic table; complex on back.
2. Assign the correct subscripts to each ion (subscript indicates the number of the ion preceding it).
Total positive charge = total negative charge.
a) Switch charges to give subscripts: charge on cation becomes subscript for anion, and vice-versa.
Use lowest whole number ratios.
OR
b) Find lowest common multiple for the charges of the 2 ions. For each ion, multiple the ion charge by a subscript number that will give the common denominator.
3. Write symbol of each ion with subscript. Do not include ion charges in final formula.
Note: For multivalent ions, the Roman numeral = charge on the cation, not the subscript.
For hydrates, prefix indicates the number of water molecules present in the ionic compound.
If more than one complex ion is present, place brackets around it, then give subscript.
Examples:
a) Monatomic: Silver chloride aluminum oxide
_________ __________
b) Polyatomic: Ammonium nitrate potassium carbonate
_________ _________
c) Multivalent: Iron(II)sulphide lead(IV)oxide
_________ _________
d) Hydrated: Zinc chloride hexahydrate Copper(II)sulphate pentahydrate
______________ ______________
Naming Ionic Compounds
General Rules:
1. Name each ion
• Monotomic: cation retains metal name, anion name ends in "ide"
• Polyatomic: give name from back of periodic table
2. Multivalent: use Roman numerals to indicate charge ( I, II, III, IV, V, VI )
3. Hydrated: use Latin prefixes to indicate # of water molecules present
1 = mono 2 = di 3 = tri 4 = tetra 5 = penta
6 = hexa 7 = hepta 8 = octa 9 = nona 10 = deca
Examples:
a) Monatomic: Ca3P2 _______________________ AlCl3 _________________________
b) Polyatomic: Na3BO3 _______________________ (NH4)2CO3 _____________________
c) Multivalent: FeO _______________________ V3N5 __________________________
d) Hydrated: BaCl2 • 2H20 _____________________________________________
Assign: Worksheets #4-7
MOLECULAR SUBSTANCES:
➢ are solids, liquids or gases at SATP
➢ if soluble, dissolve in water to form colorless aqueous solutions that do not conduct electricity
ie. they are non-electrolytes
➢ they contain only nonmetal atoms
Molecule: a particle of a molecular substance that contains a fixed number of covalently-bonded nonmetal atoms
Covalent Bond: formed from the sharing of valence electrons between nonmetal atoms, which results in an electron structure that is the same as a noble gas, for each atom in the molecule
Example: H2 A molecule of hydrogen gas has 2 atoms of Hydrogen, each with one electron. When they bond they share a pair of electrons (one pair = one covalent bond). Since each atom now has 2 electrons, they both have the same electron structure as He (noble gas).
NOTE: Ions DO NOT form molecules.
Molecular Substances Include Molecular Elements and Molecular Compounds
1. Molecular elements: contain only one kind of nonmetal atom
|Type |Molecular Elements |
| | |
|Monatomic – one atom |Noble gases He(g) Ne(g) Ar(g) Kr(g) Xe(g) Rn(g) |
| | |
|Diatomic – two atoms/molecule |Hydrogen, Oxygen, Nitrogen and the Halogens |
| | |
| |The "HONorable Halogens" |
| | |
| |H2(g) O2(g) N2(g) F2(g) Cl2(g) Br2(l) I2(s) At2(s) |
| | |
|Polyatomic – more than 2 atoms/molecule |ozone = O3(g) Phosphorus = P4(s) Sulfur (Sulphur) = S8(s) |
2. Molecular Compounds
a) Common (to memorize):
H2O(l) = water CH4(g) = methane CH3OH(l) = methanol
H2O2(l) = hydrogen peroxide C3H8(g) = propane C2H5OH(l) = ethanol
NH3(g) = ammonia C6H12O6(s) = glucose C12H22O11(s) = sucrose
b) Binary Molecular Compounds
➢ composed of 2 different kinds of nonmetals eg. CO CO2 CCl4 SO3 N2O
Writing Molecular Formulas
General Rules
1. Write each atom symbol.
2. Each prefix indicates the subscript for the nonmetal atom that precedes it (# of atoms present).
3. If no prefix is present, then there is only one atom of that nonmetal present. Monoxide = one oxygen atom present.
Examples:: Carbon monoxide _________
Carbon dioxide _________
Carbon tetrachloride _________
Sulphur trioxide _________
Dinitrogen monoxide _________
Note: Formulas for common molecular substances must be memorized, as well as those for the
"HONorable Halogens": H2(g) O2(g) N2(g) F2(g) Cl2(g) Br2(l) I2(s) At2(s)
Naming Molecular Substances
General Rules
1. First element is named in full.
2. Second element name is shortened and given an ide ending.
3. Use prefixes (same as for hydrates) to indicate the number of each kind of atom.
4. The prefix mono is usually only used for molecules with 1 atom of oxygen (monoxides).
5. Certain Hydrogen compounds (those with H first in the formula) do not use prefixes.
eg. H2S(g) = hydrogen sulfide, not dihydrogen sulfide
Examples: NO(g) ____________________________
P4O6(s) ____________________________
SO2(g) ____________________________
SO3(g) ____________________________
N2O(g) _____________________________________________
Assign: Worksheet #8
ACIDS
• Molecules that ionize in water to produce hydrogen ions, H+(aq), ions which give acids their properties
• Properties of acids:
• Conduct electricity
• Turn blue litmus paper red
• Taste sour
• React with many metals to produce hydrogen gas, H2(g)
• Have a pH value of less than 7
• Neutralize or partially neutralize bases
• General Formula: H____ (aq) or ____COOH (aq)
• Note: not all hydrogen containing compounds are acids
Eg: NH3 CH4 CH3OH C2H5OH
Naming Acids
General Rules
• Name the hydrogen compound like an ionic compound, then convert the ionic name to the acid name
hydrogen _______ide becomes hydro______ic acid
hydrogen________ite becomes __________ous acid
hydrogen________ate becomes __________ic acid
Examples:
| | | |
|Acid Formula |Ionic Name |Acid Name |
| | | |
|HCl(aq) | | |
| | | |
|HCN(aq) | | |
| | | |
|HNO2(aq) | | |
| | | |
|H2SO3(aq) | | |
| | | |
|HNO3(aq) | | |
| | | |
|H2SO4(aq) | | |
| | | |
|H3PO4(aq) | | |
| | | |
|CH3COOH(aq) | | |
Writing Acid Formulas
General Rules:
1. Translate acid name into ionic name: hydro___ic acid ( hydrogen ___ide
___ous acid ( hydrogen ___ite
___ic acid ( hydrogen ___ate
4. Write chemical formulas for each ion, using rules for writing formulas for ionic compounds.
5. Hydrogen symbol is written first (cation), except for carboxylic acids (those with COO group), in which case hydrogen is placed at the end eg: CH3COOH
6. Give the state as aqueous = (aq).
Examples:
Hydroiodic acid _____________
Chlorous acid _____________
Chloric acid _____________
Boric acid _____________
Benzoic acid _____________
Assign: Worksheet #9
BASES
• most are ionic compounds that contain the hydroxide ion, OH-, an ion that gives bases their properties
• Properties of bases:
• Conduct electricity
• Turn red litmus paper blue
• Taste bitter
• Feel slippery
• Have a pH value greater than 7
• Neutralize or partially neutralize acids
• Note: Not all compounds that contain OH are bases
Eg: CH3OH C2H5OH
Naming Bases
• follow the general rules given for ionic compounds
Examples: NaOH ____________________________
NH4OH ____________________________
Writing Base Formulas
• follow the general rules given for ionic compounds
Examples: Lithium hydroxide __________
Calcium hydroxide __________
WORKSHEET #1: ELECTRON ENERGY-LEVEL DIAGRAMS FOR ATOMS
|IA,1 | | | | | | |
| | | | | | | |
|Eg. |Sodium ion |Na+ |11 |10 |Lost 1 |1+ |
| | | | | | | |
|1 |Neon atom | | | | | |
| | | | | | | |
|2 |Lithium ion | | | |Lost 1 | |
| | | | | | | |
|3 | | |47 | | |1+ |
| | | | | | | |
|4 | | | |18 | |2- |
| | | | | | | |
|5 | |Si | | | | |
| | | | | | | |
|6 | | |33 |36 | | |
| | | | | | | |
|7 | | | |54 |Lost 1 | |
| | | | | | | |
|8 | | |30 |28 | | |
| | | | | | | |
|9 | | | |1 |0 | |
| | | | | | | |
|10 | |P | | | | |
| | | | | | | |
|11 | |Ca2+ | | | | |
| | | | | | | |
|12 |Selenide ion | | | | | |
| | | | | | | |
|13 | | |13 | | |3+ |
| | | | | | | |
|14 | |Rb+ | | | | |
| | | | | | | |
|15 | | |18 |18 | | |
| | | | | | | |
|16 | | |8 |10 | | |
| | | | | | | |
|17 |Iodine atom | | | | | |
| | | | | | | |
|18 | |Pu | | | | |
| | | | | | | |
|19 | | | |54 |Gained 2 | |
| | | | | | | |
|20 |Unnilseptium atom | | | | | |
WORKSHEET #4: NOMENCLATURE OF BINARY IONIC COMPOUNDS
| | | | | | |
|# |Chemical |Name of Compound |# |Chemical |Name of Compound |
| |Formula | | |Formula | |
| | | | | | |
|Eg. |CaCl2 |Calcium chloride |12. | |Sodium sulphide |
| | | | | | |
|1. | |Potassium chloride |13. |ZnCl2 | |
| | | | | | |
|2. |MgO | |14. | |Zinc oxide |
| | | | | | |
|3. | |Aluminum chloride |15. |KBr | |
| | | | | | |
|4. |LiBr | |16. | |Barium nitride |
| | | | | | |
|5. | |Aluminum oxide |17. |AlF3 | |
| | | | | | |
|6. |CaO | |18. | |Strontium hydride |
| | | | | | |
|7. | |Lithium nitride |19. |Cs3P | |
| | | | | | |
|8. |AgF | |20. | |Rubidium sulphide |
| | | | | | |
|9. | |Sodium chloride |21. |BeO | |
| | | | | | |
|10. |MgH2 | |22. | |Barium phosphide |
| | | | | | |
|11. | |Zinc sulphide |23. |AlBr3 | |
WORKSHEET #5: NOMENCLATURE INVOLVING POLYATOMIC (COMPLEX) IONS
| | | |
|# |Chemical Formula |Name of Compound |
| | | |
|Eg. |NaHSO4 |Sodium hydrogen sulphate |
| | | |
|1. |CaSO4 | |
| | | |
|2. |KNO3 | |
| | | |
|3. |KMnO4 | |
| | | |
|4. |Li2CO3 | |
| | | |
|5. |Ba(OH)2 | |
| | | |
|6. |NH4ClO3 | |
| | | |
|7. |Mg(NO3)2 | |
| | | |
|8. |AlBO3 | |
| | | |
|9. |Cs2SO3 | |
| | | |
|10. | |Sodium chromate |
| | | |
|11. | |Ammonium sulphide |
| | | |
|12. | |Potassium hydrogen phosphate |
| | | |
|13. | |Rubidium thiosulphate |
| | | |
|14. | |Aluminum sulphite |
| | | |
|15. | |Potassium cyanide |
| | | |
|16. | |Strontium chlorite |
| | | |
|17. | |Lithium hydroxide |
| | | |
|18. | |Ammonium silicate |
WORKSHEET #6: NOMENCLATURE INVOLVING MULTIVALENT IONS
| | | | | |
|# |Chemical |workings |Description or Use |Name of Compound |
| |Formula | |(for interest only) | |
| | | | | |
|Eg. |Cu2S | |Copper ore (chalcocite) |Copper(I)sulfide |
| | | | | |
|1. |AuCl3 | |Gold tinting of pictures | |
| | | | | |
|2. | | |Lab preparation of oxygen |Mercury(II)oxide |
| | | | | |
|3. |Sb2S3 | |Antimony ore (stibnite) | |
| | | | | |
|4. |V2O5 | |A common catalyst | |
| | | | | |
|5. | | |Forms yellow-brown solution |Iron(III)iodide |
| | | | | |
|6. | | |Copper ore (chalcopyrite) |Copper(II)sulfide |
| | | | | |
|7. |FeS | |Also in chalcopyrite | |
| | | | | |
|8. | | |Toothpaste additive |Tin(II)fluoride |
| | | | | |
|9. |MnO2 | |Managanese ore (pyrolusite) | |
| | | | | |
|10. | | |Electrode in car battery |Lead(IV)oxide |
| | | | | |
|11. | | |A green paint pigment |Chromium(III)oxide |
| | | | | |
|12. |HgS | |Mercury Ore (cinnabar) | |
| | | | | |
|13. | | |Separating types of U atoms |Uranium(VI)fluoride |
| | | | | |
|14. |SnO2 | |Tin ore (cassiterite) | |
| | | | | |
|15. | | |Uranium ore (uraninite) |Uranium(IV)oxide |
| | | | | |
|16. |Fe2O3 | |Iron ore (hematite) | |
| | | | | |
|17. | | |Forms a pink solution |Cobalt(II)chloride |
| | | | | |
|18. |TiO2 | |A white paint pigment | |
| | | | | |
|19. |NiBr2 | |Forms a green solution | |
| | | | | |
|20. | | |Forms blue solution |Copper(II)chloride |
WORKSHEET #7: NOMENCLATURE OF HYDRATED IONIC COMPOUNDS
Use Latin prefixes to indicate # of water molecules present
1 = mono 2 = di 3 = tri 4 = tetra 5 = penta
6 = hexa 7 = hepta 8 = octa 9 = nona 10 = deca
| | | |
|# |Name of Hydrate |Chemical Formula |
| | | |
|Eg. |Copper(II)sulphate pentahydrate |CuSO4 • 5H2O |
| | | |
|1 | |MgSO4 • 7H2O |
| | | |
|2 |Sodium carbonate decahydrate | |
| | | |
|3 | |MgCl2 • 6H2O |
| | | |
|4 |Barium chloride dihydrate | |
| | | |
|5 | |Cd(NO3)2 • 4H2O |
| | | |
|6 | |ZnCl2 • 6H2O |
| | | |
|7 |Zinc sulphate heptahydrate | |
| | | |
|8 |Lithium chloride tetrahydrate | |
| | | |
|9 | |Na2S2O3 • 5H2O |
| | | |
|10 |Cobalt(II)chloride hexahydrate | |
| | | |
|11 | |AlCl3 • 6H2O |
| | | |
|12 | |CaCl2 • 2H2O |
| | | |
|13 |Barium hydroxide octahydrate | |
| | | |
|14 |Nickel(II)chloride hexahydrate | |
| | | |
|15 | |Na2SO4 • 10H2O |
| | | |
|16 |Iron(III)phosphate tetrahydrate | |
| | | |
|17 | |FeSO4 • 7H2O |
| | | |
|18 |Calcium sulphate dihydrate | |
| | | |
|19 | |SnCl4 • 5H2O |
| | | |
|20 |Barium bromide tetrahydrate | |
WORKSHEET #8: NOMENCLATURE OF MOLECULAR SUBSTANCES
|# |Molecular Formula |Description or Use |Name |
| | |(for interest only) | |
| | | | |
|Eg. |CCl4 |Toxic cleaning fluid, organic solvent |Carbon tetrachloride |
| | | | |
|1 | |Component of air (78.03%) |Nitrogen |
| | | | |
|2 |O2 |Component of air (20.99%) | |
| | | | |
|3 | |Component of air (0.94%) |Argon |
| | | | |
|4 |CO2 |Component of air (0.035%) | |
| | | | |
|5 | |Component of air (0.0016%) |The other noble gases (besides Argon) |
| | | | |
|6 |NO |Air pollutant | |
| | | | |
|7 |NO2 |Air pollutant | |
| | | | |
|8 | |Air pollutant |Sulphur dioxide |
| | | | |
|9 |SO3 |Air pollutant | |
| | | | |
|10 | |Air pollutant |Carbon monoxide |
| | | | |
|11 | |Air pollutant (near surface) |Ozone |
| | | | |
|12 | |Grain alcohol, ethyl alcohol |Ethanol |
| | | | |
|13 | |Table sugar |Sucrose |
| | | | |
|14 | |Yellow solid in group VIA |Sulphur (sulfur) |
| | | | |
|15 |P4O10 |Formed from burning P4 | |
| | | | |
|16 |P4O6 |Formed from burning P4 | |
| | | | |
|17 | |Use to bleach pulp for newsprint |Chlorine dioxide |
| | | | |
|18 | |Methyl alcohol |Methanol |
| | | | |
|19 |P4 |White solid in group VA | |
| | |Used as a cleaner when dissolved in water | |
|20 | | |Ammonia |
| | | | |
|21 |CH4 |Main component of natural gas | |
| | | | |
|22 |H2O |The “universal solvent” | |
| | | | |
|23 | |Laughing gas, anaesthetic |Dinitrogen monoxide |
WORKSHEET #9: NOMENCLATURE OF ACIDS
hydrogen _______ide becomes hydro______ic acid
hydrogen________ite becomes __________ous acid
hydrogen________ate becomes __________ic acid
| | | | |
|# |Acid Formula |Workings |Acid Name |
| | | | |
|Eg. |HCl(aq) |H+ Cl- hydrogen chloride |Hydrochloric acid |
| | | | |
|1 |HBr(aq) | | |
| | | | |
|2 |H2CO3(aq) | | |
| | | | |
|3 | | |Hypochlorous acid |
| | | | |
|4 |H2CrO4(aq) | | |
| | | | |
|5 | | |Chlorous acid |
| | | | |
|6 |H2S(aq) | | |
| | | | |
|7 |H3BO3(aq) | | |
| | | | |
|8 |HI(aq) | | |
| | | | |
|9 | | |Oxalic acid |
| | | | |
|10 |HClO4(aq) | | |
| | | | |
|11 | | |Nitrous acid |
| | | | |
|12 | | |Benzoic acid |
| | | | |
|13 |H2SO3(aq) | | |
| | | | |
|14 | | |Chloric acid |
| | | | |
|15 |H2S2O3(aq) | | |
| | | | |
|16 | | |Permanganic acid |
| | | | |
|17 | | |Hydrofluoric acid |
| | | | |
|18 |HCN(aq) | | |
| | | | |
|19 | | |Thiocyanic acid |
| | | | |
|20 | | |Sulphuric acid |
SUMMARY: CLASSIFYING AND NAMING SUBSTANCES BY FORMULA
IONIC MOLECULAR
Cation + Anion All nonmetal atoms
(metal ion or NH4+) (simple nonmetal ion
or complex ion)
1. Bases: OH- = anion 1. Acids: H__(aq) or ____COOH(aq)
2. Multivalent Ionic: hydrogen ___ide ( hydro ___ic acid
- multiple charges possible on cation hydrogen ___ite ( ___ous acid
hydrogen ___ate ( ___ic acid
- use Roman numerals to indicate charge
( I, II, III, IV, V, VI )
3. Hydrated Ionic: 2. Monatomic Molecular: name element
- contain water - Noble gases
- use Latin prefixes to indicate # of water - one atom
molecules:
3. Diatomic Molecular: name element
1 = mono 6 = hexa - the "HONorable Halogens"
2 = di 7 = hepta - two atoms/molecule
3 = tri 8 = octa
4 = tetra 9 = nona 4. Common Molecular: memorize these
5 = penta 10 = deca
CH4 = methane C3H8 = propane
For all Ionic Compounds: CH3OH = methanol C2H5OH = ethanol C12H22O11 = sucrose C6H12O6 = glucose
Simple ions: cation retains metal name H2O = water P4 = phosphorus
anion name ends in "ide" H2O2 = hydrogen S8 = sulfur
peroxide O3 = ozone
Complex ions: give name from back of NH3 = ammonia
of periodic table
5. Binary Molecular:
- first element retains name
- second element ends in "ide"
- use Latin prefixes to indicate #
of each kind of atom present
- use "mono" only for oxides
SUMMARY OF PROPERTIES OF DIFFERENT SUBSTANCES
| | | | | |
|Property/ Type |Molecular |Ionic (Salts) |Acid |Base |
| | | | | |
|State |Solid (s), Liquid (l) or gas (g) |All solids (s) |Aqueous (aq) |(s) or (aq) |
| | | | | |
|Solubility |Yes or no |Yes (varies from high to |yes |Yes (varies from high to |
|in water | |low) | |low) |
| | | | | |
|Color of |Colorless |Colored or colorless |____ |_____ |
|Solution | | | | |
| | | | | |
|Conductivity |No (non-electrolyte) |Yes (electrolyte) |Yes (electrolyte) |Yes (electrolyte) |
|Of solution | | | | |
| | | | | |
|Effect on litmus paper|None |None |Blue ( Red |Red ( Blue |
SUMMARY: WRITING CHEMICAL FORMULAS
IONIC:
3. Write each ion symbol with charge: simple ions from front of periodic table; complex on back.
4. Assign the correct subscripts to each ion (subscript indicates the number of the ion preceding it).
Total positive charge = total negative charge.
b) Switch charges to give subscripts: charge on cation becomes subscript for anion, and vice-versa.
Use lowest whole number ratios.
OR
b) Find lowest common multiple for the charges of the 2 ions. For each ion, multiple the ion charge by a subscript number that will give the common multiple.
3. Write symbol of each ion with subscript. Do not include ion charges in final formula.
Note: For multivalent ions, the Roman numeral = charge on the cation, (not the subscript).
For hydrates, prefix indicates the number of water molecules present in the ionic compound.
If more than one complex ion is present, place brackets around it, then give subscript.
BASES: These are ionic hydroxides, therefore follow the rules for writing formulas for ionic compounds.
MOLECULAR:
7. Write each atom symbol.
8. Each prefix indicates the subscript for the nonmetal atom that precedes it (# of atoms present).
9. If no prefix is present, then there is only one atom of that nonmetal present. Monoxide = one oxygen atom present.
Note: Formulas for common molecular substances must be memorized, as well as those for the
"HONorable Halogens": (H2(g) O2(g) N2(g) F2(g) Cl2(g) Br2(l) I2(s) At2(s) ) and for O3, P4, & S8.
ACIDS:
1. Translate acid name into ionic name: hydro___ic acid ( hydrogen ___ide
___ous acid ( hydrogen ___ite
___ic acid ( hydrogen ___ate
2. Write chemical formulas for each ion, using rules for writing formulas for ionic compounds.
3. Hydrogen symbol is written first (cation), except for carboxylic acids (those with COO group)
10. Give the state as aqueous = (aq).
REVIEW WORKSHEET: NOMENCLATURE
Give the name or chemical formula for each of the following:
|# |NAME |FORMULA |# |NAME |FORMULA |
| | | | | | |
|1. |Calcium iodide | |26. | |SnI4 |
| | | | | | |
|2. |Silver sulphide | |27. | |HBr |
| | | | | | |
|3. |Beryllium hydride | |28. | |HBr(aq) |
| | | | | | |
|4. |Aluminum sulphate | |29. | |K2Cr2O7 |
| | | | | | |
|5. |Ammonium carbonate | |30. | |C12H22O11 |
| | | | | | |
|6. |Barium phosphide | |31. | |Ba(OH)2 |
| | | | | | |
|7. |Calcium hydrogen sulfite | |32. | |HClO4(aq) |
| | | | | | |
|8. |Sodium nitrite | |33. | |HClO2(aq) |
| | | | | | |
|9. |Magnesium fluoride | |34. | |Mg(HCO3)2 |
| | | | | | |
|10. |Nickel(II)chromate | |35. | |SF6 |
| | | | | | |
|11. |Benzoic acid | |36. | |Pb(IO3)2 |
| | | | | | |
|12. |Chromium(III)sulfite | |37. | |Mg(OH)2 |
| | | | | | |
|13. |Copper(II)nitrate tetrahydrate | |38. | |Ni3(PO4)2.(8H2O |
| | | | | | |
|14. |Bromine | |39. | |P4O6 |
| | | | | | |
|15. |Methanol | |40. | |I2 |
| | | | | | |
|16. |Ammonia | |41. | |HNO2(aq) |
| | | | | | |
|17. |Lithium chloride monohydrate | |42. | |HgCl |
| | | | | | |
|18. |Hydrofluoric acid | |43. | |HgCl2 |
| | | | | | |
|19. |Strontium hydroxide | |44. | |BaH2 |
| | | | | | |
|20. |Phosphoric acid | |45. | |AgSCN |
| | | | | | |
|21. |Aluminum nitrate | |46. | |C2H5OH |
| | | | | | |
|22. |Carbon tetrachloride | |47. | |PCl3 |
| | | | | | |
|23. |Glucose | |48. | |Na2SO3 ( 7H2O |
| | | | | | |
|24. |Lithium hydroxide | |49. | |N2O5 |
| | | | | | |
|25. |Silicic acid | |50. | |LiH2PO4 |
| | | | | | |
|51. |Nitrogen trichloride | |76. | |Ca3N2 |
| | | | | | |
|52. |Calcium sulfate dihydrate | |77. | |NaOH |
| | | | | | |
|53. |Antimony(III)oxide | |78. | |H2SO3(aq) |
| | | | | | |
|54. |Aluminum oxide hexahydrate | |79. | |H2O2 |
| | | | | | |
|55. |Magnesium benzoate | |80. | |S8 |
| | | | | | |
|56. |Dichlorine heptaoxide | |81. | |LiClO3 |
| | | | | | |
|57. |Carbonic acid | |82. | |Cd(NO3)2• 4H2O |
| | | | | | |
|58. |Phosphorous trihydride | |83. | |H3BO3(aq) |
| | | | | | |
|59. |Lead(IV)oxide | |84. | |PCl5 |
| | | | | | |
|60. |Aluminum hydroxide | |85. | |K2Cr2O7 |
| | | | | | |
|61. |Methane | |86. | |V2(CrO4)5 |
| | | | | | |
|62. |Silicon dioxide | |87. | |P2O3 |
| | | | | | |
|63. |Sodium nitrate hexahydrate | |88. | |Li3P |
| | | | | | |
|64. |Sodium acetate | |89. | |SO3 |
| | | | | | |
|65. |Oxygen | |90. | |PbSO4 |
| | | | | | |
|66. |Potassium borate | |91. | |HMnO4(aq) |
| | | | | | |
|67. |Barium phosphate | |92. | |B2H6 |
| | | | | | |
|68. |Dinitrogen trioxide | |93. | |NaHCO3 |
| | | | | | |
|69. |Acetic acid | |94. | |SnF2 |
| | | | | | |
|70. |Iron(III)phosphate tetrahydrate | |95. | |Na2S2O3 • 5H2O |
| | | | | | |
|71. |Ammonium phosphate | |96. | |HCN |
| | | | | | |
|72. |Sulfur dioxide | |97. | |HCN(aq) |
| | | | | | |
|73. |Mercury(I)oxide | |98. | |HSCN(aq) |
| | | | | | |
|74. |Sodium silicate | |99. | |SnCl4 • 5H2O |
| | | | | | |
|75. |Ozone | |100. | |NH4HSO3 |
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